A single-displacement reaction, also known as single replacement reaction or exchange reaction, is a type of chemical reaction in which one element or ligand is replaced by an atom or group. It can be represented generically as:
A + BC ⟶ AC + B {\displaystyle {\ce {A + BC -> AC + B}}}
where either
A {\displaystyle {\ce {A}}} and B {\displaystyle {\ce {B}}} are different metals (or any element that forms cation like hydrogen) and C {\displaystyle {\ce {C}}} is an anion; or
A {\displaystyle {\ce {A}}} and B {\displaystyle {\ce {B}}} are halogens and C {\displaystyle {\ce {C}}} is a cation. This will most often occur if A {\displaystyle {\ce {A}}} is more reactive than B {\displaystyle {\ce {B}}} , thus giving a more stable product. The reaction in that case is exergonic and spontaneous. In the first case, when A {\displaystyle {\ce {A}}} and B {\displaystyle {\ce {B}}} are metals, BC {\displaystyle {\ce {BC}}} and AC {\displaystyle {\ce {AC}}} are usually aqueous compounds (or very rarely in a molten state) and C {\displaystyle {\ce {C}}} is a spectator ion (i.e. remains unchanged).
A ( s ) + B + ( aq ) + C − ( aq ) ⏟ BC ( aq ) ⟶ A + ( aq ) + C − ( aq ) ⏟ AC ( aq ) + B ( s ) {\displaystyle {\ce {A(s) + \underbrace{B+(aq) + C^{-}(aq)}_{BC(aq)}-> \underbrace{A+(aq) + C^{-}(aq)}_{AC(aq)}+ B(s)}}}
In the reactivity series, the metals with the highest propensity to donate their electrons to react are listed first, followed by less reactive ones. Therefore, a metal higher on the list can displace anything below it. Here is a condensed version of the same:
K > Na > Ca > Mg > Al > C > Zn > Fe > NH 4 + > H + > Cu > Ag > Au {\displaystyle {\ce {K}}>{\ce {Na}}>{\ce {Ca}}>{\ce {Mg}}>{\ce {Al}}>{\color {gray}{\ce {C}}}>{\ce {Zn}}>{\ce {Fe}}>{\color {gray}{\ce {NH4^+}}}>{\color {gray}{\ce {H+}}}>{\ce {Cu}}>{\ce {Ag}}>{\ce {Au}}}
(Hydrogen, carbon and ammonium — labeled in gray — are not metals.) Similarly, the halogens with the highest propensity to acquire electrons are the most reactive. The activity series for halogens is:
F 2 > Cl 2 > Br 2 > I 2 {\displaystyle {\ce {F2>Cl2>Br2>I2}}}
Due to the free state nature of A {\displaystyle {\ce {A}}} and B {\displaystyle {\ce {B}}} , single displacement reactions are also redox reactions, involving the transfer of electrons from one reactant to another. When A {\displaystyle {\ce {A}}} and B {\displaystyle {\ce {B}}} are metals, A {\displaystyle {\ce {A}}} is always oxidized and B {\displaystyle {\ce {B}}} is always reduced. Since halogens prefer to gain electrons, A {\displaystyle {\ce {A}}} is reduced (from 0 {\displaystyle {\ce {0}}} to − 1 {\displaystyle {\ce {-1}}} ) and B {\displaystyle {\ce {B}}} is oxidized (from − 1 {\displaystyle {\ce {-1}}} to 0 {\displaystyle {\ce {0}}} ).
Cation replacement Here one cation replaces another:
A + BC ⟶ AC + B {\displaystyle {\ce {A + BC -> AC + B}}}
(Element A has replaced B in compound BC to become a new compound AC and the free element B.) Some examples are:
Fe + CuSO 4 ⟶ FeSO 4 + Cu {\displaystyle {\ce {Fe + CuSO4 -> FeSO4 + Cu}}}
(Blue vitriol)____(Green vitriol)
Zn + CuSO 4 ⟶ ZnSO 4 + Cu {\displaystyle {\ce {Zn + CuSO4 -> ZnSO4 + Cu}}}
(Blue vitriol)___(White vitriol)
Zn + FeSO 4 ⟶ ZnSO 4 + Fe {\displaystyle {\ce {Zn + FeSO4 -> ZnSO4 + Fe}}}
(Green vitriol) (White vitriol) These reactions are exothermic and the rise in temperature is usually in the order of the reactivity of the different metals. If the reactant in elemental form is not the more reactive metal, then no reaction will occur. Some examples of this would be the reverse.
Fe + ZnSO 4 ⟶ {\displaystyle {\ce {Fe + ZnSO4 ->}}} No Reaction
Anion replacement Here one anion replaces another:
A + CB ⟶ CA + B {\displaystyle {\ce {A + CB -> CA + B}}}
(Element A has replaced B in the compound CB to form a new compound CA and the free element B.) Some examples are:
Cl 2 + 2 NaBr ⟶ 2 NaCl + Br 2 {\displaystyle {\ce {Cl2 + 2NaBr -> 2NaCl + Br2}}}
Br 2 + 2 KI ⟶ 2 KBr + I 2 ↓ {\displaystyle {\ce {Br2 + 2KI -> 2KBr + I2(v)}}}
Cl 2 + H 2 S ⟶ 2 HCl + S ↓ {\displaystyle {\ce {Cl2 + H2S -> 2HCl + S(v)}}}
Again, the less reactive halogen cannot replace the more reactive halogen:
I 2 + 2 KBr ⟶ {\displaystyle {\ce {I2 + 2KBr ->}}} no reaction
Common reactions
Metal-acid reaction Metals react with acids to form salts and hydrogen gas.
Zn ( s ) + 2 HCl ( aq ) ⟶ ZnCl 2 ( aq ) + H 2 ↑ {\displaystyle {\ce {Zn(s) + 2HCl(aq) -> ZnCl2(aq) + H2 ^}}}
However, less reactive metals cannot displace the hydrogen from acids. (They may react with oxidizing acids though.)
Cu + HCl ⟶ {\displaystyle {\ce {Cu + HCl ->}}} No reaction
Reaction between metal and water Metals react with water to form metal oxides and hydrogen gas. The metal oxides further dissolve in water to form alkalies.
Fe ( s ) + H 2 O ( g ) ⟶ FeO ( s ) + H 2 ↑ {\displaystyle {\ce {Fe(s) + H2O (g) -> FeO(s) + H2 ^}}}
Ca ( s ) + 2 H 2 O ( l ) ⟶ Ca ( OH ) 2 ( aq ) + H 2 ↑ {\displaystyle {\ce {Ca(s) + 2H2O (l) -> Ca(OH)2(aq) + H2 ^}}}
The reaction can be extremely violent with alkali metals as the hydrogen gas catches fire. Metals like gold and silver, which are below hydrogen in the reactivity series, do not react with water.
Metal extraction Coke or more reactive metals are used to reduce metals by carbon from their metal oxides, such as in the carbothermic reaction of zinc oxide (zincite) to produce zinc metal:
ZnO + C ⟶ Zn + CO {\displaystyle {\ce {ZnO + C -> Zn + CO}}}
and the use of aluminium to produce manganese from manganese dioxide:
3 MnO 2 + 4 Al ⟶ 3 Mn + 2 Al 2 O 3 {\displaystyle {\ce {3MnO2 + 4Al -> 3Mn + 2Al2O3}}}
Such reactions are also used in extraction of boron, silicon, titanium and tungsten.
3 SiO 2 + 4 Al ⟶ 3 Si + 2 Al 2 O 3 {\displaystyle {\ce {3SiO2 + 4Al -> 3Si + 2Al2O3}}}
B 2 O 3 + 3 Mg ⟶ 2 B + 3 MgO {\displaystyle {\ce {B2O3 + 3Mg -> 2B + 3MgO}}}
TiCl 4 + 2 Mg ⟶ Ti + 2 MgCl 2 {\displaystyle {\ce {TiCl4 + 2Mg -> Ti + 2MgCl2}}}
WF 6 + 3 H 2 ⟶ W + 6 HF {\displaystyle {\ce {WF6 + 3 H2 -> W + 6 HF}}}
Thermite reaction Using highly reactive metals as reducing agents leads to exothermic reactions that melt the metal produced. This is used for welding railway tracks.
Fe 2 O 3 ( s ) + 2 Al ( s ) ⟶ 2 Fe ( l ) + Al 2 O 3 ( s ) {\displaystyle {\ce {Fe2O3(s) + 2 Al(s) -> 2 Fe(l) + Al2O3(s)}}}
a(Haematite)
3 CuO + 2 Al ⟶ 3 Cu + Al 2 O 3 {\displaystyle {\ce {3CuO + 2Al -> 3Cu + Al2O3}}}
Silver tarnish
Silver tarnishes due to the presence of hydrogen sulfide, leading to formation of silver sulfide.
4 Ag + 2 H 2 S + O 2 ⟶ 2 Ag 2 S + 2 H 2 O {\displaystyle {\ce {4Ag + 2H2S + O2 -> 2Ag2S + 2H2O}}}
3 Ag 2 S + 2 Al ⟶ 6 Ag + Al 2 S 3 {\displaystyle {\ce {3Ag2S + 2Al -> 6Ag + Al2S3}}}
Extraction of halogens Chlorine is manufactured industrially by the Deacon's process. The reaction takes place at about 400 to 450 °C in the presence of a variety of catalysts such as CuCl 2 {\displaystyle {\ce {CuCl2}}} .
4 HCl + O 2 ⟶ 2 Cl 2 + 2 H 2 O {\displaystyle {\ce {4HCl + O2 -> 2 Cl2 + 2H2O}}}
Bromine and iodine are extracted from brine by displacing with chlorine.
2 HBr + Cl 2 ⟶ 2 HCl + Br 2 ↑ {\displaystyle {\ce {2HBr + Cl2 -> 2HCl + Br2 ^}}}
2 HI + Cl 2 ⟶ 2 HCl + I 2 ↑ {\displaystyle {\ce {2HI + Cl2 -> 2HCl + I2 ^}}}
See also Double-displacement reaction Decomposition reaction Combination reaction Substitution reaction
References
External links Reactivity series by RSC Halogen displacement reaction by RSC Chlorine water reacting with Iodide and Bromide, YouTube